1. CHEMICAL QUANTITIES OR
Chapter 10
CHEMICAL QUANTITIES OR
"Our Friend the Mole"

2. How you measure how much?
You can measure mass (in grams),
or volume (in liters),
or you can count number of particles.
We count number of particles in MOLES.

3. Moles
Defined as the number of carbon atoms in exactly 12 grams of carbon-12.
1 mole is 6.02 x particles.
Treat it like a very large dozen
6.02 x is called Avogadro's number.

4. Representative particles
The smallest pieces of a substance.
For an element, it usually is an atom.
But for any diatomic element, it is a molecule
For a molecular compound it is a molecule.
For an ionic compound it is a formula unit.

5. Quick Review Molecular compounds are made of all nonmetallic elements.
CO2 NO H2O
Ionic compounds are made of ions.
All ionic compounds begin with either a metal element or NH ion.
NaCl, K2SO NH4 Br

6. What are the RP? Atom Formula unit Molecule Mg CaBr2 I2 CH4 Fe O2
NH4 NO3
NH3
Atom
Formula unit
Molecule

7. 1mol of a substance= 6.02 x 1023 RP Conversion factor
1mol NH3 = 6.02 x 1023 M.C
1mol NaCl = 6.02 x 1023 F.Unit
1mol Mg = 6.02 x 1023 Atoms

8. Example 1: How many molecules of CO2 are the in 4.56 moles of CO2 ?
Given: 4.56 moles CO2
Find: ? M.C (RP)
Conversion factor:
1mol CO2= 6.02 x M.C (RP)
Given x conversion factor:

9. Example 2: How many moles of water is 5.87 x 1022 molecules of water?
Given: 5.87 x molecules (RP)
Find: ? moles water
Conversion factor:
1mol H2 O= 6.02 x M.C (RP)
Given x conversion factor

10. Find: ? C atoms (atoms are particles but they are not RP for C6H12O6 )
Example 3: How many atoms of carbon are there in 1.23 moles of C6H12O6 ?
Given: 1.23 moles C6H12O6
Find: ? C atoms (atoms are particles but they are not RP for C6H12O6 )
Conversion factors:
Mole to M.C.(RP), then M.C (RP) to atoms.
1mol C6H12O6 = 6.02 x M.C (RP)
1 C6H12O6 M.C = 6 C atoms

11. Given x conversion factor:

12. Molar Mass
The mass of one any substance.
Aka. Gram Formula mass

13. Gram Atomic Mass
The mass of 1 mole of an element in grams.

14. Examples
How much would 2.34 moles of carbon weigh?

15. Examples
How many moles of magnesium in 4.61 g of Mg?

16. Examples
How many atoms of lithium in 1.00 g of Li?

17. Examples
How much would 3.45 x 1022 atoms of U weigh?

18. What about compounds?
in 1 mole of H2O molecules there are two moles of H atoms and 1 mole of O atoms
To find the mass of one mole of a compound
determine the moles of the elements they have
Find out how much they would weigh
add them up

19. What about compounds? What is the mass of one mole of CH4?
1 mole of C = g
4 mole of H x 1.01 g = 4.04g
1 mole CH4 = = 16.05g

20. Molar Mass The mass of 1 mole What is the molar mass of Fe2O3?
2 moles of Fe x g = g
3 moles of O x g = g
The GFM = g g = g

21. Calculate the molar mass of the following
C6H12O6
(NH4)3PO4

22. Finding moles of compounds Counting pieces by weighing
Using Molar Mass
Finding moles of compounds
Counting pieces by weighing

23. Molar Mass
The number of grams in 1 mole of atoms, formula units, or molecules.
We can make conversion factors from these.
To change grams of a compound to moles of a compound.
Or moles to grams

24. For example need to change grams to moles for NaOH
How many moles is 5.69 g of NaOH?
need to change grams to moles
for NaOH
1mole Na = 22.99g 1 mol O = g 1 mole of H = 1.01 g
1 mole NaOH = g

25. For example need to change grams to moles for NaOH
How many moles is 5.69 g of NaOH?
need to change grams to moles
for NaOH
1mole Na = 22.99g 1 mol O = g 1 mole of H = 1.01 g
1 mole NaOH = g

26. Gases and the Mole

27. Gases Many of the chemicals we deal with are gases.
They are difficult to weigh, so we’ll measure volume
Need to know how many moles of gas we have.
Two things affect the volume of a gas
Temperature and pressure
Compare at the same temp. and pressure.

28. Standard Temperature and Pressure
Avogadro's Hypothesis - at the same temperature and pressure equal volumes of gas have the same number of particles.
0ºC and 1 atmosphere pressure
Abbreviated atm
273 K and kPa
kPa is kiloPascal

29. At Standard Temperature and Pressure
abbreviated STP
At STP 1 mole of gas occupies 22.4 L
Called the molar volume
Used for conversion factors
Moles to Liter and L to mol

30. What is the volume of 4.59 mole of CO2 gas at STP?
Examples
What is the volume of 4.59 mole of CO2 gas at STP?

31. Density of a gas at STP For a gas the units will be g / L
To find the density we need the mass and the volume.
Assume you have 1 mole, then the mass is the molar mass
At STP the volume is 22.4 L.
D = m/v = molar mass/molar volume

32. Examples Find the density of CO2 at STP. D = molar mass/ molar volume
= 44.0g/ 22.4L=

33. Find the density of CH4 at STP. D = molar mass/ molar volume
Example
Find the density of CH4 at STP.
D = molar mass/ molar volume
= 16.0 g/ 22.4 L =

34. Calculating Molar Mass from Density
Use the density of a gas at STP and the molar volume at STP (22.4L/mol)
Molar mass = density (STP) x molar volume (STP)

35. Molar Mass
WHAT IS THE MOLAR MASS OF A GAS WITH A DENSITY OF G/L?
MOLAR MASS = G/L X 22.4 L /MOL = 44.0 G/MOL

36. Molar mass
What is the molar mass of a gas with a density of g/L?
Molar mass = density x molar volume
= g/ L x 22.4 L/mol
= g/ L

37. Your turn What is the density of NO2 gas at STP?
What is the molar mass of an unknown gas with a density of g/L at STP?

38. Percent Composition Part x 100 % Whole
Step 1: Find mass of each component and find total mass.
Step 2: divide mass of each component by the total mass.

39. % composition from experimental data
Calculate the percent composition of a compound that is 29.0 g of Ag with 4.30 g of S.
Mass of Ag: 29.0g
Mass of S: 4.30g
Total mass: 29.0 g g = 33.3g
Ag% = Ag/total mass = 29.0 g/ 33.3g=
S % = S/ total mass = 4.30 g/ 33.3g =

40. % composition from formula
Calculate the percent composition of C2H4?
Mass of C= 12.0 x 2 = 24.0 g
Mass of the whole compound= 24.0 g + 4x 1.01 = 28.0g
% C= 24.0 g/ 28.0 g x 100% =
% H = 4 H / C2H4 x 100% =

41. Calculate mass of each component from % composition
Mass of element A = % of element A x total mass

42. Example
How many grams of aluminum is present in 450 g of aluminum carbonate Al2 (CO3)3 ?
Step 1: find % of Al from formula
Mass of aluminum in the formula: 2 x 27g
Total formula mass: 2 x 27 + ( x 16) x 3
= g
Al % = 2 x 27 /
Step 2: Mass of Al = x 450 g =

43. How many grams of iron are there in 3.45 g of FeS?
Step 1: find % of Fe from the formula
Mass of Fe in the formula: 56 g
Total formula mass of FeS = =88 g
Fe% = 56 /88 x 100% =
Step 2: Mass of iron = % of iron x total mass

44. The Empirical and Molecular Formula
EF: The lowest whole number ratio of elements in a compound.
MF: the actual ratio of elements in a compound.
The two sometimes are the same.
MF: C 2H EF: CH2
MF: C3H EF: CH2
MF: H 2O EF: H2O

45. Interpreting EF
CH2
1 C atom combines with 2 H atom
1 dozen C atoms, 2 dozens H atoms
1 mole C atoms, 2 moles H atoms
2 moles N atoms combines with 5 moles of H atoms, what is EF?
N2 H5

46. Finding EF
4 moles of H atoms combines with 2 moles of O atoms, what is EF?
H4O2 change to H2 O
Finding EF is to find the lowest whole number mole ratio of elements in the compound

47. Calculating Empirical Formulas from % composition
Step 1: Assume you have 100 g of the compound. The percentages become grams. Step 2: Turn grams to moles using gram atomic mass of each element. Step 3: Divide all the answers by the smallest moles.

48. . 33 or.34 mole, multiply all answers by 3
Note: after step 3, if your answers are still not whole numbers, take the following steps to change your answers to whole numbers:
. 33 or.34 mole, multiply all answers by 3
.25 mole, multiply all answers by 4
. 5 mole, multiply all answers by 2

49. Calculate the empirical formula of a compound composed of 38
Calculate the empirical formula of a compound composed of % C, % H, and %N.
Step 1: Assume 100g, so we have
38,67g C, 16.22g H, 45.11g N
Step 2:
38.67 g C x 1mol C = mole C gC
16.22 g H x 1mol H = 16.1 mole H g H
45.11 g N x 1mol N = mole N gN

50. Step 3: 3.220 mol C = 1 mol C 3.220 mol EF is C1H5N1
16.1 mol H = 5 mol H mol
3.220 mol N = 1 mol N
EF is C1H5N1

51. Calculating EF
A compound is analyzed and found to contain 25.9% nitrogen and 74.1 % oxygen. What is the empirical formula of the compound?
Step 1: g N g O
Step 2: 25.9 g N x1mol/ 14 g = 1.85 mol N
Step 3: 74.1 g O x 1mol/16g = 4.63 mol O
Step 4: 1.85 mol N /1.85 = 1 mol N
4.63 mol O/ 1.85 = 2.50 mol O
Step 5: 1mol N x 2 = 2 mol N
2.5 mol O x 2 = 5 mol O
N2 O 5

52. Practice Problems
P310: Q 36, 37
P312: Q 45, 46

53. Empirical to molecular
Caffeine is 49.48% C, 5.15% H, 28.87% N and 16.49% O. What is its empirical formula?

54. Calculating MF
A compound has an empirical formula of ClCH2 and a molar mass of g/mol. What is its molecular formula?
A compound has an empirical formula of CH2O and a molar mass of g/mol. What is its molecular formula?

55. Percent to molecular Take the percent x the molar mass
This gives you mass in one mole of the compound
Change this to moles
You will get whole numbers
These are the subscripts
Caffeine is 49.48% C, 5.15% H, 28.87% N and 16.49% O. It has a molar mass of 194 g. What is its molecular formula?

56. Example
Ibuprofen is % C, 8.80 % H, % O, and has a molar mass of about 207 g/mol. What is its molecular formula?